Chapter 2 — Acids Bases And Salts

What are acids and bases?

Acids are substances that increase the concentration of hydrogen ions (H⁺) when dissolved in water; bases increase the concentration of hydroxide ions (OH⁻). This is the Arrhenius definition suitable for Class 10. Acids usually taste sour and react with metals to produce hydrogen; bases taste bitter and feel soapy.

Key properties

• Acids: sour taste, turn blue litmus red, conduct electricity in solution, react with metals to give H₂ (e.g., HCl + Zn → ZnCl₂ + H₂).
• Bases (alkalis): bitter and soapy feel, turn red litmus blue, conduct electricity in solution, react with acids to form salt and water (neutralization).

Neutralization and salts

When an acid reacts with a base, the products are usually a salt and water: acid + base → salt + water. This is called neutralization. Example: HCl + NaOH → NaCl + H₂O.

Strength vs concentration

‘Strength’ of an acid/base refers to its ability to ionize in water (strong vs weak). ‘Concentration’ refers to how much acid/base is present in a given volume (dilute vs concentrated). A strong acid can be dilute, and a weak acid can be concentrated.

pH — a measure of acidity

The acidity or basicity of a solution is expressed by pH, a logarithmic scale. pH 7 is neutral (pure water). pH < 7 is acidic; pH > 7 is basic (alkaline). The scale commonly used in Class 10 is 0–14 (at 25°C).

Importance and uses

• Acids and bases are everywhere: digestion (stomach HCl), food (vinegar, citrus), cleaning agents (bleach, baking soda), industry (fertilizers, glassmaking), and medicine (antacids).
• They are important in agriculture (soil pH affects plant growth) and the environment (acid rain).

Safety

Many acids and bases are corrosive. Use protective gear (gloves, goggles) and follow safe handling and storage rules.

Overview: In Class 10 Chemistry (Acids, Bases and Salts) the basic definitions explain what acids, bases and salts are, how they behave in water, how their strength differs from concentration, and how we detect them.

Acid (Arrhenius definition): An acid is a substance that produces hydrogen ions (H+) or hydronium ions (H3O+) when dissolved in water. Example: HCl → H+ + Cl–. Acids taste sour, turn blue litmus red, and conduct electricity in solution.

Base (Arrhenius definition): A base is a substance that produces hydroxide ions (OH–) when dissolved in water. Example: NaOH → Na+ + OH–. Bases taste bitter, feel soapy/slimy, turn red litmus blue, and conduct electricity.

Bronsted–Lowry view (brief): An acid is a proton (H+) donor and a base is a proton acceptor. This definition covers reactions in non-aqueous media as well.

Salt: A salt is an ionic compound formed when an acid reacts with a base (neutralization): acid + base → salt + water. Example: HCl + NaOH → NaCl + H2O. Salts can be neutral, acidic or basic depending on ions present.

Indicator: A substance that changes colour with pH. Natural examples: litmus (red in acid, blue in base), phenolphthalein (colourless in acid, pink in basic solutions).

Strength vs Concentration: Strength refers to extent of ionization/dissociation (strong acids/bases ionize almost completely; weak ones ionize partially). Concentration refers to amount of substance per unit volume (molarity). A dilute strong acid can be less concentrated than a concentrated weak acid but is still a strong acid chemically.

pH concept: pH is a measure of acidity/basicity: pH = −log10[H+]. pH < 7 acidic, pH = 7 neutral (at 25°C), pH > 7 basic. pH is a convenient logarithmic scale for [H+].

Electrical conductivity: Acids, bases and salts in aqueous solution produce ions and thus conduct electricity. Conductivity depends on ion concentration and mobility.

Practical notes: Many everyday substances are acids or bases: citric acid in lemon, acetic acid in vinegar, sodium bicarbonate (baking soda) is mildly basic, household ammonia (NH3 in water) is basic. Neutralization reactions are used in antacids, wastewater treatment, and agriculture (lime to neutralize acidic soil).

Definition (CBSE Class 10): Acids are substances that produce H+ (hydrogen) ions in aqueous solution. Bases are substances that produce OH− (hydroxide) ions in aqueous solution.

Basic properties:

• Acids: sour taste, turn blue litmus red, conduct electricity, react with metals to produce H2, react with bases to form salt and water (neutralisation).
• Bases: bitter taste, slippery feel, turn red litmus blue, conduct electricity, neutralise acids to give salt and water.

Strong vs weak: Strength depends on extent of ionisation (dissociation) in water. Strong acids/bases ionise almost completely (e.g., HCl, NaOH). Weak acids/bases ionise partially (e.g., CH3COOH, NH3).

Common examples with uses & simple dissociation:

• Hydrochloric acid (HCl) — strong acid. Present in stomach (gastric juice); used in labs and industry. Dissociation: HCl → H+ + Cl−.
• Sulfuric acid (H2SO4) — strong diprotic acid. Used in car batteries, fertilizers. Stepwise dissociation: H2SO4 → H+ + HSO4−; HSO4− ⇌ H+ + SO4 2−.
• Nitric acid (HNO3) — strong acid. Used to make fertilizers and explosives.
• Acetic acid (CH3COOH) — weak acid. Main component of vinegar. CH3COOH ⇌ H+ + CH3COO−.
• Carbonic acid (H2CO3) — weak acid in soft drinks (formed by CO2 + H2O).
• Sodium hydroxide (NaOH) — strong base (caustic soda). Used in soap, drain cleaners. NaOH → Na+ + OH−.
• Potassium hydroxide (KOH) — strong base, used in batteries and soaps.
• Calcium hydroxide (Ca(OH)2) — slaked lime, used in construction and to test for CO2 (turns milky).
• Ammonia (NH3) — weak base (in water: NH3 + H2O ⇌ NH4+ + OH−). Used in household cleaners and fertilizers.
• Sodium bicarbonate (NaHCO3) — mild base (baking soda); used as antacid and in baking.

Neutralisation (general): Acid + Base → Salt + Water. Example: HCl + NaOH → NaCl + H2O.

Safety note: Strong acids and bases are corrosive — handle with care, use protective gear.

Physical properties of acids, bases and salts are the observable and measurable characteristics that help distinguish them without changing their chemical identity. In everyday chemistry (CBSE Class 10), the most important physical properties are taste, touch (feel), odour, state/appearance, solubility, electrical conductivity, effect on indicators and pH.

• Taste: Acids have a sour taste (e.g., lemon, vinegar). Bases (alkalis) often taste bitter (but never taste chemicals in the lab!). Salts may be salty, sweet or tasteless depending on composition (e.g., NaCl tastes salty).
• Touch/Feel: Aqueous solutions of bases (alkalis) feel slippery or soapy because they react with fats on skin. Acids are usually corrosive and do not feel slippery.
• Odour: Many acids have a sharp or pungent smell (e.g., acetic acid vinegar). Many salts are odourless. Some bases (e.g., ammonia solution) have a distinctive pungent smell.
• State and Appearance: Most common salts are crystalline solids with high melting points. Strong acids and bases are often available as aqueous solutions; many mineral acids are liquids at room temperature (e.g., HCl solution), while some organic acids (e.g., benzoic acid) are solids.
• Solubility: Many acids, bases and salts dissolve in water to give ionic solutions. Solubility varies by compound (e.g., NaCl is highly soluble; AgCl is insoluble).
• Electrical Conductivity: When dissolved in water (or molten), acids, bases and salts that ionize conduct electricity because mobile ions carry charge. Strong electrolytes (e.g., HCl, NaOH, NaCl) conduct well; weak electrolytes (e.g., CH3COOH) conduct poorly.
• Effect on Indicators and pH: Acids turn blue litmus red and have pH < 7. Bases turn red litmus blue and have pH > 7. Neutral salts/neutral solutions have pH ≈ 7. The pH scale ranges from 0 (very acidic) to 14 (very basic).

Safety note: tastes and touches are mentioned only for conceptual identification — never taste or touch laboratory chemicals.

Definition: Acids are substances that release H+ ions (or H3O+ in water) and show characteristic chemical reactions. In Class 10 context, the main chemical properties of acids arise from the presence of hydrogen ions (H+).

Chemical properties (with observations and typical equations):

• Reaction with active metals: Acids react with metals that are above hydrogen in the reactivity series to produce a salt and hydrogen gas. Observation: effervescence; test gas with a lighted splint (gives a 'pop').
  Example equation: Zn + 2HCl → ZnCl2 + H2↑
• Reaction with metal carbonates and bicarbonates: Acids react with carbonates (CO32–) and bicarbonates (HCO3–) to give a salt, carbon dioxide gas and water. Observation: vigorous effervescence (CO2 evolution).
  Example equation: CaCO3 + 2HCl → CaCl2 + CO2↑ + H2O
• Reaction with bases (neutralization): Acids neutralize bases (metal hydroxides or oxides) to form a salt and water. These reactions are usually exothermic. Observation: temperature rise; pH moves toward 7 on neutralization.
  General equation: Acid + Base → Salt + H2O
  Example: HCl + NaOH → NaCl + H2O
• Reaction with metal oxides: Metal oxides are basic oxides; acids react with them to form salt and water.
  Example: 2HCl + CuO → CuCl2 + H2O
• Effect on indicators: Acids change the colour of indicators. Common observations: blue litmus turns red; methyl orange turns red in acidic medium; phenolphthalein is colourless in acid.
  This property is used to identify acidic solutions and to follow reactions such as titrations.

Why these reactions occur: In all above reactions the H+ ions from the acid combine with OH– (from bases) or displace hydrogen from metals, or react with carbonate species to form CO2 and H2O. The net chemical change is driven by the transfer of H+ and formation of stable products (salt, water or CO2).

pH and concentration relation (useful concept): pH measures acidity: pH = -log[H+]. Lower pH means higher [H+]. This relation explains why diluting an acid raises its pH (makes it less acidic).

Practical notes: Corrosive acids attack metals, skin and some building materials (e.g., acid rain dissolving limestone). Neutralization (with bases or antacids) is a common method to treat acid spills or excess stomach acid. Indicators and titration curves are used to determine concentration of an acid.

What is a base? A base is a substance that produces hydroxide ions (OH-) in aqueous solution (Arrhenius definition). Common bases: NaOH, KOH, Ca(OH)2, NH4OH. Bases are bitter, soapy, and feel slippery.

Chemical properties (with brief explanations)

• Action on indicators: Bases turn red litmus blue. Phenolphthalein is colourless in neutral/acidic solution and turns pink in basic solution (pH > ~8.2). Methyl orange is yellow in basic solution.
• Neutralisation with acids: Bases react with acids to form a salt and water. This is an exothermic reaction and the ionic view is H+ + OH- → H2O. Example: NaOH + HCl → NaCl + H2O.
• Reaction with ammonium salts (release of ammonia): Strong bases displace ammonia from ammonium salts. Example: NaOH + NH4Cl → NH3(g) + H2O + NaCl. The released gas has a characteristic pungent smell.
• Precipitation of metal hydroxides from salts: Bases react with solutions of many metal salts to give insoluble metal hydroxides (precipitates). Example: FeCl3 + 3NaOH → Fe(OH)3 (brown ppt) + 3NaCl.
• Reaction with acidic (non-metal) oxides / CO2: Bases react with acidic oxides such as carbon dioxide to form carbonates or bicarbonates. Example: 2NaOH + CO2 → Na2CO3 + H2O.
• Reaction with amphoteric metals/oxides/hydroxides: Some metals and their oxides/hydroxides (Al, Zn, Pb) react with strong bases to give soluble complex ions. Example: Zn(OH)2 + 2NaOH → Na2[Zn(OH)4]. Metallic aluminium reacts with NaOH producing hydrogen: 2Al + 2NaOH + 6H2O → 2Na[Al(OH)4] + 3H2 (g).
• Saponification (reaction with fats/oils): Bases like NaOH hydrolyse fats (triglycerides) to give glycerol and soap (sodium salts of fatty acids). Example (general): fat + 3NaOH → glycerol + 3 RCOONa (soap).

Important observations and uses: Neutralisation is used in medicine (antacids), industry (waste neutralisation), and household (drain cleaners). Reaction with CO2 explains why bases absorb CO2 from air (e.g., NaOH forming carbonate).

Safety note: Strong bases (e.g., NaOH, KOH) are caustic and can burn skin and eyes. Handle with care.

What is an indicator?
An indicator is a substance that changes colour depending on whether it is in an acidic or basic medium. Indicators are used to detect the presence of acids or bases and to estimate the pH of a solution.

How indicators work (chemical basis)
Most indicators are weak acids or weak bases that exist in equilibrium between two forms with different colours. For a weak acid indicator HIn:

HIn ⇌ H+ + In–

The undissociated form (HIn) has one colour and the dissociated form (In–) another. The position of the equilibrium depends on [H+]: in acidic solution the equilibrium lies to the left (more HIn), in basic solution to the right (more In–). The observed colour is the mixture of the two forms; when one form predominates the corresponding colour is seen.

Quantitative relation (useful idea)
The acid dissociation constant for the indicator (Ka) is

Ka = [H+][In–] / [HIn]

Taking logs and rearranging gives the Henderson–Hasselbalch type relation for the indicator:

pH = pKa + log([In–]/[HIn])

When [In–] = [HIn] (ratio 1:1) the pH = pKa and both colours are equally visible. Practically, a visible colour change occurs roughly when pH = pKa ± 1 (this defines the indicator's transition range).

Types of indicators

• Natural indicators — obtained from plants (e.g., litmus from lichens, red cabbage extract, turmeric).
• Synthetic indicators — prepared in labs (e.g., phenolphthalein, methyl orange, methyl red, universal indicator).

Choosing an indicator
Choose an indicator whose transition range includes the pH at the equivalence point of the reaction you are monitoring. For example, for titrations that have an equivalence pH in the basic region, use phenolphthalein; for acidic-equivalence titrations, use methyl orange.

Common indicators and their approximate transition ranges

• Litmus: ~pH 4.5–8.3 (acid: red; base: blue)
• Methyl orange: ~pH 3.1–4.4 (acid: red; base: yellow)
• Methyl red: ~pH 4.4–6.2 (acid: red; base: yellow)
• Phenolphthalein: ~pH 8.2–10.0 (acid: colourless; base: pink)
• Universal indicator: shows a range of colours across pH 1–14

Behaviour examples (practical notes)
- In titrations the rapid pH change at equivalence is matched with an indicator whose colour change lies within that steep region.
- Natural indicators (e.g., red cabbage) give different colours across a wide pH range and are useful for demonstration or rough pH estimation.
- Indicators are not used to give exact pH values; for precise pH use a pH meter.

Limitations

• Indicator colour depends on concentration and medium; very dilute solutions may not show colours well.
• Some indicators are affected by temperature or redox conditions.

Summary: An indicator changes colour because its acid and base forms have different colours. The equilibrium between these forms depends on pH; using Ka and pKa (Henderson–Hasselbalch) we can predict the pH range over which an indicator will change colour and thus select the appropriate indicator for experiments and practical checks.

What is pH? pH is a measure of the acidity or basicity of an aqueous solution. It is defined as the negative base‑10 logarithm of the hydrogen‑ion (H+) concentration: pH = -log10[H+]. The pH scale commonly runs from 0 to 14 at 25°C: values below 7 are acidic, 7 is neutral, and above 7 are basic (alkaline).

Why logarithmic? Because pH uses a logarithm, each unit change in pH corresponds to a tenfold change in H+ concentration. For example, a solution at pH 3 has ten times more H+ than one at pH 4 and 100 times more than one at pH 5.

Key features

• Neutral water (at 25°C): [H+] = [OH-] = 1×10^-7 M so pH = 7.
• Relation with hydroxide ions: pOH = -log[OH-], and at 25°C pH + pOH = 14 because Kw = [H+][OH-] = 1×10^-14.
• Temperature dependence: Kw (and thus the numerical value 14) changes with temperature, so neutral pH is 7 only at 25°C.
• Strong acids/bases: for a strong monoprotic acid (e.g., HCl) that is fully dissociated, [H+] ≈ initial molarity.
• Weak acids/bases: need equilibrium (Ka or Kb) to find [H+] or [OH-]; approximate formulas are used when dissociation is small.

pH in real systems — pH affects biological systems, corrosion, chemical reactions and industrial processes. Living organisms maintain narrow pH ranges (e.g., human blood ≈ 7.35–7.45). Indicators such as litmus, phenolphthalein and methyl orange change colour at characteristic pH ranges and help estimate pH.

Quick practical notes

• To calculate pH of a strong acid: if [H+] = c (M), pH = -log c.
• For a weak acid HA with concentration c and acid constant Ka (assuming x << c): [H+] ≈ sqrt(Ka·c) and pH ≈ -log(sqrt(Ka·c)). Equivalently pH ≈ 1/2 (pKa - log c).
• Buffers resist pH change by containing a weak acid and its conjugate base (or weak base and its conjugate acid).

Definition — Strength: 'Strength' of an acid or base refers to its ability to ionize (dissociate) in water. A strong acid/base ionizes almost completely (e.g., HCl, NaOH); a weak acid/base ionizes only partially (e.g., CH3COOH, NH3). Strength is an intrinsic chemical property and is quantified by the acid dissociation constant Ka (or base dissociation constant Kb).

Definition — Concentration: 'Concentration' refers to how much acid or base is present in a given volume of solution. Common measures: molarity (c, mol L−1), percent (w/v or w/w), ppm. Concentration is an extrinsic property and can be changed by dilution or mixing.

How they differ: Strength tells how completely molecules ionize; concentration tells how many molecules are present per unit volume. A strong acid can be very dilute (low concentration) but remain 'strong' (still fully ionizes proportionally), and a weak acid can be very concentrated but remain 'weak' (still only partially ionizes).

Effect on [H+] and pH:

• For a strong monoprotic acid at concentration c (mol L−1), [H+] ≈ c and pH = −log10 c.
• For a weak acid HA with dissociation HA ⇌ H+ + A−, Ka = [H+][A−]/[HA]. If initial concentration is c and degree of dissociation is α, then [H+] = cα, [A−] = cα, [HA] ≈ c(1−α).

Relation between degree of dissociation (α), Ka and concentration:

Ka = (cα)(cα)/(c(1−α)) = c α^2/(1−α). For small α (typical of weak acids) 1−α ≈ 1, giving the useful approximation α ≈ sqrt(Ka / c). This shows α increases as c decreases (i.e., dilution increases percent ionization of weak acids).

Key consequence: On dilution, weak acids/bases show increased degree of ionization (α↑) and hence larger fraction of molecules ionized — the solution becomes relatively more conductive and pH shifts more than expected from linear dilution. Strong acids/bases are essentially completely ionized at all ordinary concentrations, so their degree of ionization is independent of concentration; however, the absolute [H+] (and thus pH) still changes with concentration.

Common student confusion: People often confuse 'dilute HCl' (low concentration) with 'weak HCl' (HCl is a strong acid regardless of concentration). Remember: strength = intrinsic ionization tendency; concentration = amount present.

Definition: A neutralisation reaction is a chemical reaction in which an acid and a base react to form a salt and usually water. In terms of ions, H+ (from the acid) combines with OH- (from the base) to form H2O. Neutralisation reactions are usually exothermic.

General chemical form:

• Acid + Base → Salt + Water
• Ionic (net) equation: H+ (aq) + OH- (aq) → H2O (l)

Details & important points:

• When a strong acid reacts with a strong base (both fully dissociated), the equivalence point (true neutralisation) occurs near pH 7 and about one mole of H+ combines with one mole of OH- to give one mole of water.
• Neutralisation is exothermic. For strong acid + strong base the standard heat of neutralisation is about −57 kJ per mole of water formed.
• Not all neutralisation reactions produce only water and a neutral salt; some acids react with carbonates/bicarbonates to give salt + CO2 + H2O (e.g., NaHCO3 + HCl → NaCl + CO2 + H2O).
• Salts formed can be neutral, acidic, or basic depending on whether the parent acid/base were strong or weak (e.g., NaCl is neutral; NH4Cl is acidic; Na2CO3 is basic).
• Neutralisation is used in titrations to determine unknown concentrations by using an indicator (or pH meter) to find the end point (close to equivalence point). Common indicators: litmus, phenolphthalein, methyl orange. The choice depends on the type of acid/base being titrated.

Connection to titration: During a titration, plot of pH vs volume of titrant gives a characteristic curve. The steepest region corresponds to the equivalence point. For strong acid + strong base the curve is steep and the equivalence pH ≈ 7. For weak acid + strong base the equivalence pH > 7; for strong acid + weak base the equivalence pH < 7.

Practical significance: Neutralisation reactions are used in medicine (antacids), agriculture (lime to neutralise acidic soil), environmental protection (treating acidic industrial effluents, neutralising acid spills), food (baking soda neutralising excess acidity), and household cleaning.

What is a salt? A salt is an ionic compound formed when the H+ of an acid is replaced by a metal ion or an ammonium ion (NH4+). Salts are generally formed by neutralisation, displacement or double‑displacement (precipitation) reactions.

Main ways salts are formed

• Neutralisation (acid + base): An acid reacts with a base (metal hydroxide) to give a salt and water. Example: HCl + NaOH → NaCl + H2O. Ionic view: H+ + OH− → H2O.
• Acid + metal: A metal above hydrogen in the reactivity series reacts with an acid to give a salt and hydrogen gas. Example: Zn + 2HCl → ZnCl2 + H2↑.
• Acid + metal oxide: Metal oxides (basic oxides) react with acids to give salt and water. Example: CuO + 2HCl → CuCl2 + H2O.
• Acid + metal carbonate / bicarbonate: Produces salt, carbon dioxide and water. Example: CaCO3 + 2HCl → CaCl2 + CO2↑ + H2O.
• Double displacement / precipitation (salt + salt): Two soluble salts react to form an insoluble salt (precipitate) and another soluble salt. Example: AgNO3 + NaCl → AgCl(s)↓ + NaNO3.
• Acid + ammonium compounds: Acid reacts with ammonium salts to liberate ammonia or form ammonium salts. Example: HCl + NH3 → NH4Cl.

Types of salts (based on neutralisation)

• Normal (neutral) salts: Complete neutralisation; solution pH ≈ 7 (e.g., NaCl).
• Acidic salts: Incomplete neutralisation of a polyprotic acid (e.g., NaHSO4 from partial neutralisation of H2SO4).
• Basic salts: Incomplete neutralisation of polyvalent base or decomposition of hydroxide (e.g., Cu2(OH)3Cl or basic lead carbonate).

pH of salt solutions: The pH of a salt solution depends on the ions produced and their hydrolysis. Examples: NaCl (neutral), NH4Cl (acidic because NH4+ ⇌ NH3 + H+), Na2CO3 (basic because CO32− + H2O ⇌ HCO3− + OH−).

Hydrated salts: Many salts crystallise with water of crystallisation (eg. CuSO4·5H2O, Na2CO3·10H2O). Heating can remove water and change properties.

Everyday relevance: Salts are common in food (NaCl), medicine (NaCl IV fluids; NH4Cl in expectorants), agriculture (NPK fertilizers such as KNO3, NH4NO3), construction (gypsum—CaSO4·2H2O → wall plaster), water treatment (salts in hard water), and household products (washing soda Na2CO3).

Overview
A salt is an ionic compound formed when the hydrogen ion of an acid is wholly or partly replaced by a metal or another positive ion. In the laboratory salts are prepared by different chemical reactions (neutralisation, displacement, carbonate reaction and precipitation) and are purified by techniques such as filtration, evaporation, crystallisation and recrystallisation.

Common laboratory methods of preparation

• 1. Neutralisation (acid + base): A soluble salt is made by reacting an acid with a base (metal hydroxide or alkali). The reaction gives salt and water. Example: HCl + NaOH → NaCl + H₂O. To obtain solid crystals, the solution is evaporated until saturated and then allowed to cool so crystals form (crystallisation).
• 2. Acid + Metal Carbonate/Bicarbonate: Acid reacts with a carbonate to give salt, CO₂ and water. Example: 2HCl + Na₂CO₃ → 2NaCl + CO₂ + H₂O. After CO₂ evolution stops, filter if necessary and crystallise the filtrate.
• 3. Acid + Metal: A metal (more reactive than hydrogen) reacts with acid to give a salt and hydrogen gas. Example: 2HCl + Zn → ZnCl₂ + H₂. The resulting solution is evaporated/crystallised to get solid salt.
• 4. Precipitation (double displacement): Mix two soluble salt solutions to produce an insoluble salt that precipitates out. Example: AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq). The precipitate is collected by filtration, washed and dried.

Purification techniques

• Filtration — separates insoluble impurities or precipitates from liquid. Use gravity filtration (filter paper + funnel) or vacuum filtration (Buchner funnel) for faster drying.
• Washing — the precipitate is washed with cold distilled water to remove soluble impurities (mother liquor).
• Evaporation and Crystallisation — concentrate the solution by gentle heating to approach saturation, then cool to allow crystals to form. Collect crystals by filtration and dry them.
• Recrystallisation — dissolve impure solid in minimum hot solvent, filter hot to remove insoluble impurities (hot filtration), then cool so pure crystals form. This improves purity by exploiting different solubilities.
• Drying — place filtered crystals/precipitate on a watch glass or in a desiccator or oven (low temperature) to remove physically-adsorbed water without decomposing the salt.
• Fractional crystallisation — separates two salts in a mixture using their different solubilities (cooling a concentrated solution so the less-soluble salt crystallises first).

Typical laboratory procedure (example: preparing copper(II) sulfate crystals)

1. React copper oxide (CuO) with dilute sulphuric acid until in excess of one reactant is used: CuO + H₂SO₄ → CuSO₄ + H₂O.
2. Filter to remove unreacted insoluble CuO.
3. Evaporate the filtrate to concentrate; when hot and saturated, allow to cool to form blue crystals of CuSO₄·5H₂O.
4. Collect crystals by filtration, wash with small amount of cold water, and dry.

Practical tips & safety

• Use gentle heat to avoid splattering and decomposition. Do not boil to dryness unless the salt is stable on heating.
• Perform reactions producing gases (H₂, CO₂) in a well-ventilated area or under a fume hood.
• Use appropriate personal protective equipment (gloves, goggles, lab coat).

Key terms: mother liquor (the leftover solution after crystallisation), filtrate (liquid passed through filter), residue (solid left on filter), saturation point, recrystallisation, precipitate.

What is a salt?
A salt is an ionic compound formed when an acid reacts with a base (neutralisation) or by replacement of H+ in an acid by a metal or a positive radical. Example: HCl + NaOH → NaCl + H2O.

Classification of salts (Class 10 level)

• Neutral salts: Formed from a strong acid and a strong base. Their aqueous solutions are neutral (pH ≈ 7) because neither ion hydrolyses significantly. Example: NaCl (from HCl + NaOH).
• Acid salts (acidic salts): Formed by partial neutralisation of a polyprotic acid (or when a strong acid reacts with a weak base). They still contain replaceable H+ and give acidic solutions. Examples: NaHCO3 (sodium hydrogen carbonate, from partial neutralisation of H2CO3) and NaHSO4 (from partial neutralisation of H2SO4). Another type is salt from strong acid + weak base: NH4Cl (from HCl + NH3) — solution is acidic because NH4+ hydrolyses to produce H+.
• Basic salts (alkaline salts): Formed by partial neutralisation of a polyhydroxy base or when a weak acid reacts with a strong base. Their solutions are basic because anion hydrolyses to give OH−. Example: NaCH3COO (sodium acetate, from CH3COOH + NaOH) — acetate ion hydrolyses to give OH−.
• Double salts: Crystalline salts that contain two different simple salts in their lattice and show combined properties in the solid state but dissociate into constituent ions in solution. They are formed by mixing solutions of two salts and crystallising. Examples: Mohr's salt FeSO4·(NH4)2SO4·6H2O, Potash alum KAl(SO4)2·12H2O. Note: Double salts are different from complex (coordination) compounds; double salts dissociate into simple ions in solution.
• Normal salts: (sometimes called 'basic/common' salts) — salts in which all replaceable H+ of the acid has been replaced by metal (e.g., NaCl, K2SO4). This term is often used to contrast with acid/basic salts.

Why solutions are acidic or basic — hydrolysis
When salts dissolve, their ions can react (hydrolyse) with water. The tendency to hydrolyse depends on the strengths of the parent acid and base:

• Salt of strong acid + strong base: neither ion hydrolyses → neutral solution (pH ≈ 7).
• Salt of strong acid + weak base: cation hydrolyses (produces H+) → acidic solution (pH < 7). Example: NH4Cl → NH4+ + Cl−; NH4+ + H2O ⇌ NH3 + H+.
• Salt of weak acid + strong base: anion hydrolyses (produces OH−) → basic solution (pH > 7). Example: CH3COONa → CH3COO− + Na+; CH3COO− + H2O ⇌ CH3COOH + OH−.
• Salt of weak acid + weak base: pH depends on relative strengths (pKa and pKb) of the parent acid and base.

Identification (simple tests)
Use litmus: acidic salts turn blue litmus red; basic salts turn red litmus blue; neutral salts do not change litmus. Observe pH with indicator paper or pH meter for precise measurement.

Typical classroom examples and observations
- Add NaHCO3 (baking soda) to water: slightly basic (CO3(2−)/HCO3− hydrolysis).
- Dissolve KAl(SO4)2·12H2O (alum): neutral to slightly acidic depending on hydrolysis of Al3+ ions.
- Mohr's salt forms stable, bright crystals used in volumetric analysis.

Important distinctions
- Acid/basic salt refers to partial neutralisation or hydrolysis behavior in water.
- Double salts exist as unique crystals but dissociate into ions in solution; complex salts (coordination compounds) retain complex ions in solution.

Teaching tip: Use a simple flowchart: start with 'which parent acid/base are strong or weak?' → classify salt (neutral/acidic/basic). Show examples and litmus test results alongside.

Definition: Hydrated salts are crystalline ionic compounds that contain a definite number of water molecules chemically bound in their crystal lattice. The water present in the crystal is called the water of crystallization (or crystallization water). A general formula is MxAy·nH2O, where n is the number of water molecules per formula unit.

Nature and properties:

• The number of water molecules (n) is fixed for a particular hydrated salt (e.g., CuSO4·5H2O has n = 5).
• Hydrated salts often have characteristic colours caused or modified by the coordinated water (e.g., blue CuSO4·5H2O; when dehydrated it becomes pale/white).
• On heating, hydrated salts may lose some or all of their water of crystallization (dehydration). Some dehydration reactions are reversible — the anhydrous salt can reabsorb water (rehydration).
• Related phenomena: efflorescence (loss of water to atmosphere), deliquescence (absorb moisture and dissolve), hygroscopicity (absorb moisture but not necessarily dissolve).

How to determine water of crystallization experimentally (typical lab method):

1. Weigh a clean, dry crucible with the hydrated salt sample: mass = m1.
2. Heat strongly until all water is lost; cool in a desiccator and weigh the crucible with anhydrous salt: mass = m2.
3. Mass of water lost = m1 - m2. Mass of anhydrous salt = m2 - mass of empty crucible.
4. Convert masses to moles: n(H2O) = mass of water / 18.0 g mol⁻¹; n(salt) = mass of anhydrous salt / M (molar mass of anhydrous part).
5. Number of water molecules per formula unit n = n(H2O) / n(salt). Round to nearest whole number.

Sample calculation (CuSO4·nH2O): Suppose a sample of copper(II) sulfate hydrate weighed 2.500 g (including crucible) and after heating the anhydrous salt weighed 1.636 g (same crucible correction already applied). Mass of water lost = 2.500 − 1.636 = 0.864 g. Molar masses: H2O ≈ 18.0 g mol⁻¹; CuSO4 ≈ 159.6 g mol⁻¹.

moles H2O = 0.864 / 18.0 = 0.0480 mol. moles CuSO4 = 1.636 / 159.6 = 0.01025 mol. n ≈ 0.0480 / 0.01025 ≈ 4.69 ≈ 5. So the hydrate is CuSO4·5H2O.

Chemical behaviour (examples of dehydration/rehydration):

• CuSO4·5H2O (blue) → CuSO4 (white/gray) + 5 H2O (g) (on heating)
• CoCl2·6H2O (pink) → CoCl2 (blue) + 6 H2O (g)
• MgSO4·7H2O (Epsom salt) → MgSO4 + 7 H2O
• CaSO4·2H2O (gypsum) heated partially → CaSO4·½H2O (plaster of Paris) + 3/2 H2O

Importance and applications: Hydrated salts are important in drying, manufacture, construction and laboratory work. Examples include washing soda (Na2CO3·10H2O), Epsom salt (MgSO4·7H2O) in medicine, gypsum (CaSO4·2H2O) in plaster, and copper sulfate in agriculture. Some salts are used as humidity indicators (cobalt(II) chloride).

Safety and practical notes: Heating must be controlled — some salts decompose rather than simply lose water; use a desiccator to prevent reabsorption before weighing.

Summary: Water of crystallization is fixed and part of the crystal structure. It can be removed by heating and its quantity determined experimentally by mass loss. Hydrated salts show characteristic physical properties (colour, solubility) that change on dehydration.

What is salt hydrolysis?
Salt hydrolysis is the reaction of ions from a dissolved salt with water to produce H₃O⁺ or OH⁻, which changes the pH of the solution. Whether a salt solution is acidic, basic or neutral depends on the strengths of the parent acid and base.

Classification (simple rules)

• Salt of strong acid + strong base → neutral solution (pH ≈ 7). Example: NaCl.
• Salt of strong acid + weak base → acidic solution. Example: NH₄Cl (NH₄⁺ hydrolyses to give H₃O⁺).
• Salt of weak acid + strong base → basic solution. Example: CH₃COONa (A⁻ hydrolyses to give OH⁻).
• Salt of weak acid + weak base → pH depends on relative strengths; compare K_a and K_b of the conjugates.

Basic hydrolysis reactions (examples)

• CH₃COO⁻ + H₂O ⇌ CH₃COOH + OH⁻ (gives basic solution)
• NH₄⁺ + H₂O ⇌ NH₃ + H₃O⁺ (gives acidic solution)

Equilibrium constants and relations
At 25 °C, K_w = [H⁺][OH⁻] = 1.0×10⁻¹⁴ and pK_w = 14.

• For an anion A⁻ (conjugate base of weak acid HA): K_b (for A⁻) = K_w/K_a(HA).
• For a cation BH⁺ (conjugate acid of weak base B): K_a(BH⁺) = K_w/K_b(B).

Approximate pH formulas (for a salt concentration c, mol L⁻¹)

• Salt of weak acid (A⁻) from strong base: A⁻ hydrolyses, K_b = K_w/K_a. [OH⁻] ≈ √(K_b · c).
  pOH ≈ 1/2 (pK_b − log c).
  pH = 14 − pOH.
• Salt of weak base (BH⁺) from strong acid: BH⁺ hydrolyses. [H⁺] ≈ √(K_a · c) where K_a(BH⁺) = K_w/K_b.
  pH ≈ 1/2 (pK_a − log c).
• Percent (degree) of hydrolysis for A⁻: α ≈ √(K_b/c). Percent hydrolysis ≈ 100·α = 100·√(K_b/c). (Derived from K_b = x²/(c − x) ≈ x²/c when x ≪ c.)
• Salt of weak acid + weak base (both conjugates present): if both hydrolyses occur, an approximate result often used for class 10 is pH ≈ 7 + 1/2 (pK_a − pK_b). Interpretation: if pK_a > pK_b the solution is basic; if pK_a < pK_b acidic; if equal ≈ 7.

How the formulas are obtained (short derivation)
For A⁻: A⁻ + H₂O ⇌ HA + OH⁻, K_b = [HA][OH⁻]/[A⁻]. If initial salt conc. = c and x = [OH⁻] formed, K_b ≈ x²/c ⇒ x ≈ √(K_bc).

Limitations and approximations
These simple formulas assume x ≪ c (small degree of hydrolysis) and temperature 25 °C (K_w = 1×10⁻¹⁴). For very dilute solutions or very strong hydrolysis (large K_b or K_a) you must solve the full equilibrium expressions.

Common observations

• Diluting a basic salt (e.g., CH₃COONa) lowers its pH toward 7 (absolute [OH⁻] decreases), though degree of hydrolysis increases.
• Metal ions with high positive charge (Al³⁺, Fe³⁺) hydrolyse appreciably and give acidic solutions.

Acids, bases and salts are widely used in everyday life, industry, agriculture and medicine because of their characteristic chemical behaviours — acidity/basicity, reactivity with metals and carbonates, neutralisation and salt formation. Understanding these applications explains why certain substances are chosen for specific tasks and shows how neutralisation is routinely used to control pH.

Household uses

• Cleaning — basic solutions (e.g., caustic soda, NaOH) dissolve grease and unclog drains; weak acids (vinegar, citric acid) remove limescale and rust.
• Cooking and food preservation — acetic acid (vinegar) and citric acid are food preservatives and flavouring agents; common salt (NaCl) is used for seasoning and preservation.
• Personal care — soaps and detergents are produced using NaOH (saponification); baking soda (NaHCO3) is used in baking and as a mild antacid.

Industrial and manufacturing uses

• Sulfuric acid (H2SO4) is a major industrial chemical used in fertilizer manufacture (superphosphates), petroleum refining, battery acid and chemical synthesis.
• Hydrochloric acid (HCl) is used for pickling steel, removing rust, and in the production of chloride salts.
• Chlor-alkali process produces NaOH, Cl2 and H2; NaOH is used in paper, soap and textile industries.
• Nitric acid (HNO3) is used to make fertilizers and explosives; hydrofluoric acid (HF) etches glass (hazardous).

Agriculture and environment

• Fertilizers are salts containing nitrogen, phosphorus or potassium (e.g., NH4NO3, KCl, Ca(H2PO4)2) that supply nutrients to plants.
• Soil pH correction uses lime (CaO or Ca(OH)2) to neutralise acidic soils; neutralisation reactions are used to treat acidic industrial effluents.
• Bleaching powder (CaOCl2) and chlorine-containing salts are used for disinfection.

Medicine and safety

• Antacids (containing Mg(OH)2, Al(OH)3 or NaHCO3) neutralise excess stomach acid to relieve acidity/heartburn.
• pH control in pharmaceuticals is critical for drug stability and activity; many medicinal salts deliver active ions in safe forms.

Laboratory and analytical uses

• Titration (acid–base titration) is used to determine concentration of acids or bases using indicators or pH meters. The neutralisation point (equivalence point) is a key concept.
• Indicators (litmus, methyl orange, phenolphthalein) are dyes that change colour over specific pH ranges and are used to detect acidity/basicity.
• Electroplating and electrolysis use salt solutions as electrolytes to deposit metals or produce industrial chemicals.

Safety note: Many concentrated acids and bases are corrosive and hazardous. Neutralisation with the appropriate reagent, dilution and use of protective equipment are essential when handling them.

Summary: The characteristic reactions of acids, bases and salts — neutralisation, reaction with metals and carbonates, and ion exchange — make them indispensable across household, industrial, agricultural and medical applications. Practical use often involves controlling pH, choosing suitable salts and applying neutralisation reactions safely.

The chapter topic "Environmental and Safety Aspects" deals with how acids, bases and salts affect the environment and human safety, and what practical measures are taken to reduce harm. This includes pollution caused by acidic or basic discharges, acid rain, effects on soil and aquatic life, corrosion of buildings, methods of neutralization and treatment of effluents, and laboratory/industrial safety practices.

Environmental effects

• Acid rain: Sulphur dioxide (SO2) and nitrogen oxides (NOx) from burning fossil fuels are converted in the atmosphere into strong acids (SO3 → H2SO4, NO2 → HNO3) and fall as acid rain. Typical reactions: SO2 + 1/2 O2 → SO3; SO3 + H2O → H2SO4. Acid rain lowers pH of lakes and soils causing plant damage, killing aquatic organisms and mobilizing toxic metal ions (e.g., Al3+) in soil.
• Water bodies: Discharge of acidic or alkaline industrial effluents changes pH, harming aquatic flora and fauna. Sudden pH shifts can denature enzymes and disrupt biological processes.
• Soil and agriculture: Acidic soils decrease nutrient availability and increase toxic metal solubility. Farmers often add lime (CaCO3 or Ca(OH)2) to neutralize soil acidity.
• Corrosion and building damage: Acid rain reacts with carbonate rocks (e.g., limestone): H2SO4 + CaCO3 → CaSO4 + CO2 + H2O, causing erosion and structural damage.
• Eutrophication (salts): Excessive nitrates and phosphates (salts) from fertilizers lead to algal blooms, reducing dissolved oxygen and killing aquatic life.

Safety aspects (laboratory, school, industry)

• Storage: Store strong acids and bases separately in labeled, corrosion-resistant containers and cabinets. Keep incompatible chemicals apart.
• PPE and equipment: Use goggles, gloves, lab coat, and face shield when handling corrosives. Have eyewash stations and safety showers accessible.
• Handling & dilution: Always add acid to water (not water to acid) to dilute strong acids—this minimizes heat splashes. Similarly, add concentrated base slowly to water.
• Spill response: For small spills, neutralize acids with a weak base (e.g., sodium bicarbonate) and neutralize bases with a weak acid (e.g., dilute acetic acid) only if trained. Absorb neutralized residue and dispose as instructed. For skin/eye contact, rinse immediately with plenty of water and seek medical attention.
• Waste treatment: Industrial acidic/alkaline effluents are typically neutralized (e.g., with lime Ca(OH)2) and treated to remove harmful ions before discharge. Proper effluent treatment reduces environmental impact.

Practical preventive and remedial measures

• Limit SO2/NOx emissions via scrubbers, catalytic converters and cleaner fuels to reduce acid rain.
• Neutralize acidic soils by liming; monitor soil pH regularly.
• Treat wastewater: neutralization, precipitation of heavy metals, biological treatment for organic pollutants.
• Use buffer solutions where pH must be maintained (e.g., aquaria, industrial processes).

Key takeaways: pH control is central to protecting ecosystems and human health. Safe handling, correct neutralization and proper waste treatment prevent damage from acids, bases and salts.

Overview
"Tests and Observations" covers simple qualitative and quantitative procedures used to identify acids, bases and their reactions. Common observations include colour changes with indicators, effervescence with carbonates or metals, slippery feel for bases, and heat release on neutralization.

Indicator tests
Indicators are substances that change colour depending on pH. Typical observations:

• Blue litmus paper turns red in an acid; red litmus turns blue in a base.
• Methyl orange: red in acid, yellow in alkaline medium.
• Phenolphthalein: colourless in acid and neutral, pink/magenta in basic solutions (around pH > 8.2).

Reactions giving characteristic observations

• Acid + metal (e.g., Zn, Mg) → effervescence due to H2 gas. Test H2 with a burning splint (produces a ‘pop’ sound).
• Acid + carbonate/bicarbonate → effervescence of CO2. Confirm CO2 by passing gas through limewater (Ca(OH)2) which turns milky (forms CaCO3).
• Acid + base → neutralization: salt + water; often exothermic (temperature rises). If an indicator is present, it shows endpoint colour change.
• Bases often feel slippery (due to saponification of skin oils) and neutralize acids; they give pink colour with phenolphthalein.

Qualitative observations to note in the laboratory

• Colour of indicator before and after adding sample.
• Effervescence (rate/intensity) and behaviour of gas (test identity).
• Temperature change on mixing acid and base (use thermometer).
• Solubility/precipitate formation (e.g., CaCO3 white precipitate in limewater).

Precautions & interpretation
Indicator colours depend on concentration and type of indicator. Some acids/bases are weak and may give smaller changes. Always perform gas tests safely and use dilute samples when possible.

Overview
This topic covers the key chemical reactions and equations from the chapter "Acids, Bases and Salts" (Class 10). Emphasis is on neutralisation, reactions of acids and bases with metals, metal carbonates, metal oxides and ammonium salts, salt hydrolysis, and a few important preparative and real-life reactions.

1. Neutralisation (acid + base)
General: acid + base → salt + water. Neutralisation removes H⁺ from acid and OH^- from base to form H₂O.

Examples (balanced):
HCl + NaOH → NaCl + H₂O
H₂SO₄ + 2 NaOH → Na₂SO₄ + 2 H₂O

2. Acids with metals
Acids react with many metals (above hydrogen in activity series) to give a salt and hydrogen gas.

Example:
Zn + 2 HCl → ZnCl₂ + H₂

3. Acids with metal carbonates / bicarbonates
Acid + carbonate/bicarbonate → salt + CO₂ + H₂O. This is the evolution of carbon dioxide gas (effervescence).

Examples:
CaCO₃ + 2 HCl → CaCl₂ + CO₂ + H₂O
2 NaHCO₃ + H₂SO₄ → Na₂SO₄ + 2 CO₂ + 2 H₂O

4. Acids with metal oxides (basic oxides)
Acid + metal oxide → salt + water.

Example:
2 HCl + CuO → CuCl₂ + H₂O

5. Bases with ammonium salts (release of NH₃)
Base + ammonium salt → salt + ammonia + water (ammonia gas is liberated on heating or on strong base addition).

Example:
NH₄Cl + NaOH → NaCl + NH₃ + H₂O

6. Decomposition of bicarbonates on heating
Many bicarbonates decompose on heating giving carbonate, water and CO₂.

Example:
2 NaHCO₃ (s) → Na₂CO₃ (s) + H₂O (g) + CO₂ (g)

7. Saponification (bases react with fats)
Triglyceride (fat/oil) + NaOH → glycerol + sodium salts of fatty acids (soap).

Typical equation (schematic):
Fat + 3 NaOH → Glycerol + 3 RCOONa (soap)

8. Preparation of salts
Methods include neutralisation (acid + base), precipitation (mixing two soluble salts to form insoluble salt), and direct reactions (metal + acid, metal oxide + acid). Example of precipitation:
AgNO₃ + NaCl → AgCl (s) + NaNO₃

9. Salt hydrolysis (effect of salt on pH)
Salts from strong acid + strong base are neutral (e.g., NaCl). Salts from strong acid + weak base are acidic (e.g., NH₄Cl), and salts from weak acid + strong base are basic (e.g., Na₂CO₃). Hydrolysis reactions:

NH₄⁺ + H₂O → NH₃ + H₃O⁺ (acidic)
CO₃^2- + H₂O → HCO₃^- + OH^- (basic)

10. Important observational reactions
- Lime water test for CO₂: CO₂ + Ca(OH)₂ → CaCO₃ (white precipitate)
- Ammonium chloride with NaOH gives pungent NH₃ smell and turns red litmus blue.
- Phenolphthalein: colourless in acid, pink in base. Methyl orange: red in acid, yellow in base.

Key concepts to remember
- Stoichiometry: balance H⁺ and OH^- for neutralisation.
- Reaction types: acid + metal, acid + carbonate, acid + oxide, base + ammonium salt, saponification, precipitation.
- Behaviour of salts in water depends on the strength of parent acid/base (hydrolysis).